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*STUDENT*
*STUDENT*
Regents Chemistry
NOTE PACKET
Unit 1: Matter & Measurement
Copyright © 2015 Tim Dolgos
Copyright © 2015 Tim Dolgos
Copyright © 2015 Tim Dolgos
*STUDENT*
*STUDENT*
Unit Vocabulary:
1.
2.
3.
4.
5.
6.
7.
8.
9.
S.I. unit
Meter
Liter
Gram
Mass
Weight
Volume
Density
Intensive
10.
11.
12.
13.
14.
15.
16.
17.
18.
Extensive
Significant Figures
Precision
Accuracy
Matter
Element
Compound
Mixture
Heterogeneous Mixture
19.
20.
21.
22.
23.
24.
25.
Homogeneous Mixture
Pure Substance
Particle Diagram
Chromatography
Filtration
Distillation
Scientific Notation
Unit Objectives: When you complete this unit you will be able to do the following…
1) Classify types of matter
2) Draw particle diagrams to represent different types of matter
3) Recognize various techniques that can be used to separate matter
4) Convert between units of measurements
5) Differentiate between accuracy and precision
6) Write numbers in scientific notation
7) State rules to determine significant figures
8) Count significant figures
9) Understand the importance of significant figures
10) Calculate the volume and density of an object
Copyright © 2015 Tim Dolgos
Matter
Can NOT be separated
by physical means
Can NOT be separated
by chemical means
Particle
Diagram
CAN be Separated by
PHYSICAL means
Same
composition
throughout
Separated by chemical
means, only
Particle
Diagram
Particle
Diagram
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Different
composition
throughout
Particle
Diagram
Copyright © 2015 Tim Dolgos
Practice Problems:
1. Which particle diagram(s) represent a mixture?
2. Which particle diagram(s) represent a pure substance?
3. Which of the following particle diagrams represents a mixture of
one compound and one element?
4. Which particle diagram represents a diatomic element?
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Copyright © 2015 Tim Dolgos
Properties of Matter:
 Physical properties are the constants about a substance; can use
our senses to observe them; do not require chemical analysis
Example:
o Extensive Property: a property that depends on how
much material you are dealing with
Ex:
o Intensive Property: a property that does not depend on
how much material you are dealing with (help identify
matter; a constant about that particular type of matter)
Ex:
 Chemical properties include behaviors substances adhere to when they
__________ with other substances
Examples:
Guided Practice: Identify the following as being intensive, extensive, or chemical
properties.
____________ 1. The mass of copper wire is 255 g.
____________ 2. The boiling point of ethyl alcohol is 77°C.
___________ 3. Baking soda reacts with vinegar to make carbon dioxide gas.
____________ 4. The density of mercury is 13.6g/mL.
____________ 5. The solubility of sodium chloride in water is 40g/100mL of water.
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Copyright © 2015 Tim Dolgos
Physical vs. Chemical Changes
Matter is always changing. Ice in your drink melts. Wood in
your fire burns.
Physical Change – a change that does NOT alter the chemical properties
of a substance (example: ______________, ______________); change in
size or shape; ____________________________
*PHYSICAL processes can be reversed ________________________
Example: ice melting to become liquid (its still water!)
Chemical Change – a reaction in which the composition of a substance is
changed (example: __________); properties _______________________
1. Signs of a chemical rxn:
1.
2.
3.
Example: firewood burning
Change of Matter
Physical or Chemical?
Burning toast
Making ice cubes
Lighting a candle
Spoiling milk
Making kool-aid
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Copyright © 2015 Tim Dolgos
Elements vs. Compounds
Element = _____________________________________________________
Compound = ____________________________________________________
1. Circle ( ) all the elements and underline the compounds below.
2. On the line provided, record the number of different symbols within the species.
CO
___
C2H5OH
Mg
___
H2SO4 ___
O2
___
C
___
___
Al(CN)3 ___
He
___
___
Cl2
___
NI3
H2O ___
Cu
Co
___
___
NaCl
___
I
___
Questions:
1) Does each compound have the same number of symbols? ____
2) For each ELEMENT above, how many total symbols are listed? __
3) What is the minimum number of symbols that must be present in
order for a species to be considered a compound? __
Understanding Compound Formulas:
 Within a compound, you may see subscripts. These subscripts tell you the
number of each type of atom that is present.
Example:
# carbon atoms __
# oxygen atoms __
 If there are parentheses present around two or more atoms, the subscript
applies to all atoms within the parentheses.
Example:
# aluminum atoms __
# carbon atoms __
# nitrogen atoms __
 If one of the atoms within the parentheses has a subscript, you multiply this
number by the number outside of the parentheses.
Example:
# iron atoms __
# sulfur atoms __
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# oxygen atoms ___
Copyright © 2015 Tim Dolgos
There is no vodcast for this page. Just know that you will be responsible for memorizing
the symbol and name for the most commonly used elements, which appear below. You will
be quizzed on each of the two sets below at some point next week. I recommend making
flash cards for all of them (there’s a flashcard app, and my former students loved it) and
spend a couple minutes a night going over them.
The Common Elements
Rules for writing element symbols:
1) ______________________________________
2) ______________________________________
* Symbol *
Ag
Al
Ar
As
Au
B
Ba
Be
Br
C
Ca
Cl
Co
Cr
Cs
Cu
F
Fe
Fr
H
He
Hg
* Name *
silver
aluminum
argon
arsenic
gold
boron
barium
beryllium
bromine
carbon
calcium
chlorine
cobalt
chromium
cesium
copper
fluorine
iron
francium
hydrogen
helium
mercury
* Symbol *
I
K
Kr
Li
Mg
Mn
N
Na
Ne
Ni
O
P
Pb
Ra
Rb
Rn
S
Si
Sn
Sr
U
Xe
Zn
* Name *
iodine
potassium
krypton
lithium
magnesium
manganese
nitrogen
sodium
neon
nickel
oxygen
phosphorus
lead
radium
rubidium
radon
sulfur
silicon
tin
strontium
uranium
xenon
zinc
MEMORIZE both directions (symbol to name, name to symbol) for Quiz on _____________
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Copyright © 2015 Tim Dolgos
Separation of Matter
Separation Apparatus
Type of
Separation
(Physical or
Chemical)
Description of
Technique
What types of
matter will it
separate?
Filtration
Watch Glass Evaporation
Crucible Evaporation
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Copyright © 2015 Tim Dolgos
Separation of Matter
(continued)
Distillation
Chromatography
On the other hand  _____________________ requires
reacting a sample with something else in order to turn it into a
completely different compound
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SCIENTIFIC NOTATION –
method for expressing very large or
small numbers easily (Example: ___________________)
For example, the number 1,000,000 is in standard formation format.
The scientific notation of this number is 1.0 x 106



We always move the decimal place to make the ____________(the number
out in front) between _______________
We then arrange the ___________ (the number up to the right of the ten)
Now, _______________  if you were to take the 1.0 and move the
decimal place 6 places to the right (since it is a positive number), you would
get the original number (1,000,000)
Example: 123000000000
Guided Practice – Write the following numbers in scientific notation
(remember the mantissa rule!)
1. 34000000 =
2. 0.0000067 =
3. 25,864 =
Now, write the following scientific notations in standard (normal)
notation form:
4. 5.7 x 108 =
5. 6.34 x 10-11 =
Calculator Practice:
First, let’s enter the number 2.3 x 10-5 in scientific notation:
1. Type “2”
2. Type the decimal point
3. Type “3”
4. Then press the “ee” “EXP” or “
” key(s)
5. Press the “+/-“ key (NOT the “—“ or “subtract” key)
6. Type “5”
Next, let’s enter that number by 1 mole, or 6.02 x 1023. What do you get for
your answer? _________________
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Copyright © 2015 Tim Dolgos
Measurements and the Metric System
In chemistry we measure matter using ____ units. This is an abbreviation for
_________________________________.
SI BASE UNITS (AKA Base Units):
**If you forget, use Table D in your Reference Tables!
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Copyright © 2015 Tim Dolgos
SI Metric Prefixes
Prefix Symbol
tera
T
giga
G
mega
M
kilo
k
hecto
h
deca
da
no prefix:
deci
d
centi
c
milli
m
micro

nano
n
pico
p
femto
f
atto
a
Numerical (Multiply Root Word
by)*
1,000,000,000,000
1,000,000,000
1,000,000
1,000
100
10
1
0.1
0.01
0.001
0.000001
0.000000001
0.000000000001
0.000000000000001
0.000000000000000001
Exponential
1012
109
106
103
102
101
100
10¯1
10¯2
10¯3
10¯6
10¯9
10¯12
10¯15
10¯18
*Example: In the word kilometer, the root word (base unit) is “meter” and
the prefix is “kilo.” Kilo means multiply the root word by 1000. Therefore,
one kilometer is 1000 meters (1 km = 1000 m).
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Copyright © 2015 Tim Dolgos
Conversion Factors – a mathematical expression that relates two units that
measure the same type of quantity
Examples:
-
*Rest Assured! For the Regents, the most you will have to convert will be between the
milli-/kilo-/base unit (g, L, etc.). This is always a matter of ___________________.
You must also make sure you move the decimal the ___________________ (right or
left, which depends on whether you are converting from small to big or vice versa).
TRICK:
kilo
hecto
deca
base unit
deci
centi
milli
k
h
d
base unit
d
c
m
Let’s practice!
1. A car travels 845 km. How many meters is this?
2. Convert 0.0290 L to milliliters.
3. Convert 2500mL to liters.
9. 12 mL = ______ L
4. 3 g = _______ kg
5. 1 km = ______ m
Compare by placing a <, >, or = on
the line provided:
6. 1 kg = _______ g
10.
56 cm
11.
7g
7. 1 L = ________ mL
8. 7 m = _______ mm
__ 6 m
__ 698 mg
Once you get your answer, check it! Does it make sense?
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Copyright © 2015 Tim Dolgos
Dimensional Analysis
Often you will be required to solve a problem with mixed units, or to convert from one set
of units to another. Dimensional analysis is a simple method to accomplish this task.
Ex: How many minutes are there in 15 days?
Solution A:
STEP 1: Figure out the units that you have and the steps to get to the units
that you need.
HAVE
(What’s missing?)
NEED
Days (d)
Minutes (min)
STEP 2: Make a “grid” and plug in the numbers to make your first conversion.
The number/units you HAVE goes in the top left, the number/units
you NEED go in the top right, and the conversion factor goes in the
bottom right.
Need
Have
15 d 24 h
1d
=
Conversion Factor
STEP 3: Cancel “like terms.” Then, multiply the top numbers (the numerators)
together and divide the result by the bottom number (the
denominator).
15 d 24 h
1d
= 360 h
Since 24 hours and 1 day are equivalent, you are actually multiplying
15 days by a factor of 1. This means that the magnitude of your
number stays the same and only the units change.
In other words, 15 days = 360 hours
STEP 4: Now, use your answer from Step 3 as the new “HAVE” and repeat
the process using the conversion factor 60 minutes = 1 hour
360 h 60 min
1h
=
21,600 min
Now you try on: How many minutes are there in the month of October?
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Copyright © 2015 Tim Dolgos
ACCURACY VS. PRECISION
Accuracy – ____________________________________________
Ex: Hitting bulls eye when you are aiming for it
*For most experiments, ______________ means _____ from
the expected value
Precision – _____________________________________________
_____________________________________________________
*For an experiment with +/- 5% as the margin for accuracy,
that means the difference between the highest and lowest
percent error cannot exceed a ____________________
Ex: If the highest percent error for an experiment is
+7.6%, and the lowest is -5.4% that range is 13.0%,
which means that experiment was not precise
Practice: Cheryl, Cynthia, Carmen, and Casey take target practice in PE. Assuming
that they were all aiming at the bulls eye, match each target with the proper
description.
(a) Accurate and precise
(b) Not precise, but one piece of data accurate
(c) Precise but not accurate
(d) Neither precise nor accurate
Practice: The following data was collected during a lab experiment. The density of
the cube should be 10.8 g/mL. Is this data is accurate, precise, both, or neither?
Justify your answer. ______________________________________________
______________________________________________________________
Trial Number
1
2
3
Density of Cube
6.2 g/mL
6.3 g/mL
6.5 g/mL
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Copyright © 2015 Tim Dolgos
SIGNIFICANT FIGURES
- also known as Sig Figs (SF)
 A method for handling ____________________ in all measurements
 This arises due to the fact that we have different equipment with different
degrees of ___________________
 Significant figures are associated with ____________________________
 _______________________ do _______________ when determining sig figs
o Ex: Atomic masses on periodic table
Conversions (1in = 2.54 cm)
Examples:
1. Reading a ruler

We know for sure that the object is more than _____, but less than _____

We know for sure that the object is more than _____, but less than _____

This ruler allows us to estimate the length to ________
2. Reading a graduated cylinder:
► Measurements are read from the bottom of the _________
►Which gives a volume reading of _______
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Copyright © 2015 Tim Dolgos
The Atlantic/Pacific Method - another way to determine the # sig figs in a number
1)
2)
Determine if a decimal point is present. If a decimal is present, think of “P”
for “present.” If there is no decimal, think of “A” for “Absent.” P stands
for the Pacific coast and A stands for the Atlantic Coast.
Imagine the number you are looking at is a map of the USA. Begin counting
from the correct side of the number (Atlantic/right side or Pacific/left
side) based on what you determined in step 1. Consider the first nonzero
number you land on the start of your count. Consider each digit from here
on out significant as well until you reach the other end of the number.
Pacific Coast
3.
Atlantic Coast
Decimal is
Present
Decimal is
Absent
1. Start @ 1st
NONZERO
1. Start @ 1st
NONZERO
2. Count all
the way to the
Atlantic—NO
EXCEPTIONS
2. Count all
the way to the
Pacific—NO
EXCEPTIONS
Determine the number of significant numbers in each of the following:
1) 357
_______
5) 0.0357
2) 3570
_______
6) 3.570 x 103
_______
3) 3570.
_______
7) 0.3570
_______
4) 0.357
_______
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_______
Copyright © 2015 Tim Dolgos
Rules for Determining Number of Significant Figures in a Given Number
Rule
Example
1. All nonzero numbers (ex: 1 – 9) are always 123456789 m
significant
1.23 x 102
2. Zeros located between nonzero numbers
are significant
40.7 L
87,009 km
3. For numbers less than one, all zeros to
the left of the 1st nonzero number are
NOT significant
0.009587 m
4. Zeros at the end of a number and to the
right of a decimal point are significant
85.00 g
5. Zeros at the end of a whole number may
be significant or not. If there is a decimal
after the last zero, they are significant.
If there is not a decimal point after the
end zeros, they are NOT significant
6. Exact numbers have an infinite
number of significant figures
2000 m
PRACTICE:
Measurement
1020 mL
1200 m
1200. L
1200.00 mm
0.001 km
10.00 L
12000 m
00.100 cL
22.101 mm
101,000 km
0.0009 kg
9.070000000 L
2000. m
1 ft = 12 inch
Number of Significant Figures
20
Rule(s) Applied
Copyright © 2015 Tim Dolgos
Rules for Using Sig Figs in Calculations
General Rule  Final answer must be expressed in the lowest amount of significant figures
that were originally given to you (you can’t create accuracy when you didn’t have it to start
with!)
Operation
Rule
Multiplication/Division
Perform operation as
normal & express
answer in least # sig
figs that were given to
you
Addition/Subtraction
Examples:
Examples
Line decimal points up;
round final answer to
lowest decimal place
(least accurate) value
given
12.257 x 1.162 =
+
3.95
2.879
213.6____
5.1456 – 2.31 = _______
69.25/45.8 = _________
Rules for Calculations with Numbers in Scientific Notation:
Rule
Example
Addition/Subtraction  All values must 4.5 x 106 - 2.3 x 105
have the same exponent. Result is the
sum or difference of the mantissas,
multiplied by the same exponent of 10
Multiplication  mantissas are
multiplied and exponents of 10 are
(3.1 x 103) (5.01 x 104)
added
Division  mantissas are divided and
exponents are subtracted
7.63 x 103 / 8.6203 x 104
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Copyright © 2015 Tim Dolgos
MEASURING MATTER
1. Mass vs. Weight
MASS
WEIGHT
*We really only work with ________ in chemistry class!
** We have the same _________ whether we are on earth or on the moon.
The different forces of gravity on each cause us to weigh more on earth
than on the moon though (this is why we float on the moon!)
2. Volume - amount of _____________ an object takes up
 Techniques:
Liquids 
Regular Solids 
Irregular Solids 
3. Density: amount of mass in a given space; _________ of mass to volume
Formula (Table T):
BOX A
BOX B
Which box has a higher density? Explain your answer.
____________________________________________________
____________________________________________________
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Copyright © 2015 Tim Dolgos
Density Problems – Show all work!
*Note: the density of water is ______________
1) What is the density of an object with a mass of 60 g and a volume of 2 cm3?
2) If you have a gold brick that is 2.0 cm x 3.0 cm x 4.0 cm and has a mass of
48.0 g, what is its density?
3) If a block of wood has a density of 0.6 g/ cm3 and a mass of 120 g, what is
its volume?
4) What is the mass of an object that has a volume of 34 cm3 and a density of
6.0 g/cm3?
5) Which is heavier, a ton of feathers or a ton of bowling balls?
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Copyright © 2015 Tim Dolgos
Percent Error

Measurement of the % that the measured value is “off” from accepted value
Measured value =
Accepted value =

Formula is found in Table T (back page 12) of your Reference Tables:
If negative, your measured value is ________________ the accepted value
If positive, your measured value is ________________ the accepted value
*It is very important that you put the given values into the proper place in the
formula!
Sample Problem: In a lab experiment, you are told by your teacher that the actual
(or accepted) amount of sugar in a can of Coke is 39 g. You experimentally
determine it to be 40 g based on your own data and calculations. What is your
percent error? Express answer in the proper amount of significant figures.
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Copyright © 2015 Tim Dolgos